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Unit 1 – Thermochemistry Review
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There are two major outcomes in this unit:
1. Determine and interpret energy changes in chemical reactions
2. Explain and communicate energy changes in chemical reactions
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Key Concepts:
 enthalpy of formation
 enthalpy of reaction
 H notation
 Hess’ law
 molar enthalpy
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energy diagrams
activation energy
catalysts
calorimetry
fuels and energy efficiency
Specific Outcomes:
 Apply Q = m c t to the analysis of heat transfer
 Explain, in a general way, how stored energy in the chemical bonds of
hydrocarbons originated from the Sun
 Define enthalpy and molar enthalpy for chemical reactions
 Write balanced equations for chemical reactions that include energy changes
 Use and interpret H notation to communicate and calculate energy changes in
chemical reactions
 Predict the enthalpy change for chemical equations using standard enthalpies of
formation
 Explain and use Hess’ law to calculate energy changes for a net reaction from a
series of reactions
 Use calorimetry data to determine the enthalpy changes in chemical reactions
 Identify that H2O(l) and CO2(g) are reactants in photosynthesis and products of
cellular respiration while H2O(g) and CO2(g) are the products of hydrocarbon
combustion in an open system
 Classify chemical reactions as endothermic or exothermic, including those for the
processes of photosynthesis, cellular respiration and hydrocarbon combustion
 Design a method to compare the molar enthalpy change when burning two or
more fuels, identifying and controlling major variables
 Compare energy changes associated with a variety of chemical reactions through
the analysis of data and energy diagrams
 Define activation energy as the energy barrier that must be overcome for a
chemical reaction to occur
 Explain the energy changes that occur during chemical reactions, referring to
bonds breaking and forming and changes in EP and EK
 Analyze and label energy diagrams of a chemical reaction, including reactants,
products, enthalpy change, and activation energy
 Explain that catalysts increase reaction rates by providing alternate energy
pathways for changes without affecting the net amount of energy involved
 Draw enthalpy diagrams, indicating changes in energy for chemical reactions
 Explain the discrepancy between the theoretical and actual efficiency of a thermal
energy conversion system
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Unit A Review Questions
1. If beaker A contains 400 mL of a liquid and beaker B contains 800 mL of the same
liquid, what happens when 250 kJ of energy is absorbed by the liquid in each beaker?
2. Classify each of the following as endothermic or exothermic:
a)
b)
c)
d)
e)
boiling water
forming water
photosynthesis
combustion of methane
cellular respiration
3. “The reactants have more potential energy than the products do.” What kind of
reaction does this statement describe? Justify your answer.
4. You are given a 50 g sample of each of the metals shown below, all at 25oC. You
heat each metal under identical conditions. Which metal will be first to reach 30oC?
Which will be last? Explain your reasoning.
Metal
platinum
titanium
zinc
Specific Heat Capacity (J/goC)
0.133
0.523
0.388
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5. Draw and label a potential energy diagram for the following chemical reaction:
A(s) + B(g)  AB(s) + 397 kJ
Ea(fwd) = 27 kJ
Include activation energy (forward and reverse), the transition state, and the enthalpy
change for the reaction.
6. Add the following labels to the diagram you drew for question 5:
 bonds being broken
 bonds being formed
 kinetic energy  potential energy
 potential energy  kinetic energy
7. Write balanced thermochemical equations for the formation of the following
substances at standard conditions:
a) CH3COOH(l)
b) BaSO4(s)
c) C2H2(g)
8. Rank the substances in question 7 from least thermally stable to most thermally
stable.
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9. Sulfur hexafluoride, SF6(g), is one of the densest known gases and is used to insulate
electrical equipment, including high voltage wires. Determine the enthalpy change of
reaction (1) below using reactions (2), (3), and (4).
(1)
H2S(g) + 4 F2(g)  2 HF(g) + SF6(g)
H = ?
(2)
½ H2(g) + ½ F2(g)  HF(g)
H = - 273 kJ
(3)
(4)
S8(s) + 24 F2(g)  8 SF6 (g)
8 H2(g) + S8(s)  8 H2S(g)
H = - 9760 kJ
H = - 168 kJ
10. Nitroglycerine, C3H5(NO3)3(l), an explosive, detonates to produce nitrogen, water
vapour, carbon dioxide, and oxygen gases, releasing 1.15 x 104 kJ/mol of
nitroglycerine.
a. Write a balanced chemical equation for the reaction.
b. Using standard molar enthalpies of formation, determine the molar enthalpy of
formation of nitroglycerine.
11. n-propanol has a chemical formula of C3H7OH(l) and a fH of -302.6 kJ/mol. When
16.5 g of n-propanol is burned, it causes the temperature of 5.00 L of water to
increase by 19.5oC. Calculate the efficiency of this process.